Experimental Procedure

1. Experiment 6. Galvanic Cells, the Nernst Equation Experimental Procedure

2. Overview The cell potentials for a number of galvanic cells are measured and the redox couples are placed in order of decreasing reduction potentials. The effects of changes in ion concentrations on cell potentials are observed and analyzed. Perform the experiment with a partner. At each circled superscript [1-12] in the procedure, stop and record your observation on the Report Sheet. Discuss your observation with your lab partner and your instructor.

3. A. Reduction Potentials of Several Redox Couples Collect the electrodes, solutions, and equipment. 1) Obtain four small (~50 mL) beaker and fill them three-fourths full of the 0.1 M solutions as shown in Figure 32.3. Share these solutions with other groups of chemists in the laboratory. 2) Polish strips of copper, zinc, magnesium, and iron metal with sand paper, rinse briefly with dilute (~1 M) HNO3 (Caution!), and rinse with deionized water.

4. 3) The polished metals, used as electrodes, should be bent to extend over tip of their respective beakers. Check out a multimeter (Figure 32.4) with two electrical wires attached to alligator clips. Figure 32.3 Setup for measuring the cell potentials of six galvanic cells

5. 2. Set up the copper-zinc cell. Place a Cu strip (electrode) in the CuSO4 solution and a Zn strip (electrode) in the Zn(NO3)2 solution. Roll and flatten a piece of filter paper; wet the filter paper with a 0.1 M KNO3 solution. Fold and insert the ends of the filter paper into the solutions in two beakers; this is the salt bridge shown in picture next page. Set the multimeter to the 2000-mV range or as appropriate. Connect one electrode to the negative terminal of the multimeter and the other to the positive terminal.

7. 3. Determine the copper-zinc cell potential. If the multimeter reads a negative potential, reverse the connections to the electrodes. Read and record the (positive) cell potential. Identify the metal strips that serve as the cathode (positive terminal) and the anode. Write an equation for the half-reaction occurring at each electrode. Combine the two half-reactions to write the equation for the cell reaction.[1]

8. 4. Repeat for the remaining cells. Determine the cell potentials for all possible galvanic cells that can be constructed from the four redox couples. Refer to the Report Sheet for the various galvanic cells. Prepare a new salt bridge for each galvanic cell.[2]

11. 5. Determine the relative reduction potentials. Assuming the reduction potential of the Zn2+(0.1 M)/Zn redox couple is -0.79 V, determine the reduction potentials of all other redox couples. [3] 6. Determine the reduction potential of the unknown redox couple. Place a 0.1 M solution and electrode obtained from your TA in a small beaker. Determine the reduction potential, relative to the Zn2+ (0.1 M)/Zn redox couple, for your redox couple.[4]

13. B. Effect of Concentration Changes on Cell Potential Effect of different molar concentrations. Set up the galvanic cell shown in Figure 32.5, using 1 M CuSO4 and 0.001 M CuSO4 solutions. Immerse a polished copper electrode in each solution. Prepare a salt bridge (Part A.2) to connect the two half-cells. Measure the cell potential. Determine the anode and the cathode. Write an equation for the reaction occurring at each electrode.[5]

14. 2. Effect of complex formation. Add 2-5 mL of 6 M NH3 to the 0.001 M CuSO4 solution until any precipitate redissolves. (Caution: Do not inhale NH3.) Observe and record any changes in the half-cell and the cell potential.[6] 3. Effect of precipitate formation. 1) Add 2-5 mL of 0.2 M Na2S to the 0.001 M CuSO4 solution now containing the added NH3. What is observed in the half-cell and what happens to the cell potential? Record your observations.[7]

15. Figure 32.5 Setup for measuring the cell potentials of a Cu2+ concentration cell

16. C. The Nernst Equation and an Unknown Concentration Prepare the diluted solutions. Prepare solutions 1 through 4 as shown in Figure 32.6 using a 1-mL pipet and 100-mL volumetric flasks. Be sure to rinse the pipet with the more concentrated solution before making the transfer. Use deionized water for dilution to the mark in the volumetric flasks. Calculate the molar concentration of the Cu2+ ion for each solution and record.[8]

17. Figure 32.6 Successive quantitative dilution, starting with 0.1 M CuSO4

18. 2. Measure and calculate the cell potential for solution 4. Set up the experiment as shown in Figure 32.7, page 356, using small (~50 mL) beakers. The Zn2+ /Zn redox couple is the reference half-cell for this part of the experiment. Connect the two half-cells with a new salt bridge. Reset the multimeter to the lowest range (~200 mV). Connect the electrodes to the multimeter and record the potential difference, Ecell,expt.[9] Calculate the theoretical cell potential Ecell,expt. (Use a table of standard reduction potentials and the Nernst equation.)[10]

19. Figure 32.7 Setup to measure the effect that diluted solutions have on cell potentials

21. 3. Measure and calculate the cell potential for solutions 3 and 2. Repeat Part C.2 with solutions 3 and 2, respectively. A freshly prepared salt bridge is required for each cell. 4. Plot the date. Plot Ecell,expt and Ecell,calc (ordinate) versus pCu (abscissa) on the same piece of linear graph paper (page 362) or by using appropriate software for the four concentrations of CuSO4 (see data from Part A.3 for the potential of solution 1). Have your TA approve your graph. [11]

22. 5. Determine the concentration of the unknown. Obtain a CuSO4 solution with an unknown copper ion concentration from your TA and set up a like galvanic cell. Determine Ecell as in Part C.2. Using the graph, determine the unknown copper(II) ion concentration in the solution.[12]

23. Students should thoroughly clean the electrodes, multimeter, and all connectors before returning this equipment.

24. Cleanup Rinse the beakers twice with tap water and twice with deionized water. Discard the rinses in the Waste Metal Solutions container.